Chemistry · Ch 6 — Chemical Bonding and Molecular Structure
Molecular Orbital Theory
Molecular Orbital Theory
Molecular Orbital Theory
The molecular orbital (MO) theory was developed by F. Hund and R.S. Mulliken in 1932. It offers a fundamentally different picture of chemical bonding from the valence bond approach. Instead of imagining bonds formed by overlapping atomic orbitals that remain localised between two atoms, MO theory treats the entire molecule as a single quantum system where electrons occupy orbitals that belong to the whole molecule.
The Core Idea: Electrons in Molecules
The central premise is straightforward: just as electrons in an atom occupy atomic orbitals (1s, 2s, 2p, etc.), electrons in a molecule occupy molecular orbitals. A molecular orbital is a wave function that describes the probability distribution of an electron in the field of all the nuclei in the molecule. The key difference is that an atomic orbital is monocentric (influenced by one nucleus), while a molecular orbital is polycentric (influenced by two or more nuclei).
How Molecular Orbitals Form
Molecular orbitals are created by the linear combination of atomic orbitals (LCAO). For this combination to be effective, the atomic orbitals must satisfy two conditions:
- Comparable energies — orbitals that are very different in energy (like a 1s and a 2p) do not combine effectively.
- Proper symmetry — the orbitals must overlap in a way that their wave functions can constructively or destructively interfere.
When two atomic orbitals combine, they always produce exactly two molecular orbitals. This is a direct consequence of quantum mechanics: the number of molecular orbitals formed equals the number of atomic orbitals that combine.
If atomic orbitals combine, they produce exactly molecular orbitals. For two atomic orbitals, we get two molecular orbitals.
Bonding and Antibonding Molecular Orbitals
Of the two molecular orbitals formed, one is called a bonding molecular orbital and the other an antibonding molecular orbital.
- The bonding molecular orbital has lower energy than the original atomic orbitals. It is formed by constructive interference of the atomic wave functions, which increases electron density between the nuclei. This stabilises the molecule.
- The antibonding molecular orbital has higher energy than the original atomic orbitals. It is formed by destructive interference, which creates a node (region of zero electron density) between the nuclei. This destabilises the molecule.
The bonding orbital is always more stable (lower in energy) than the corresponding antibonding orbital. This energy difference is what drives chemical bonding — electrons preferentially occupy the lower-energy bonding orbitals.
Filling Molecular Orbitals with Electrons
The rules for filling molecular orbitals are identical to those for atomic orbitals:
- Aufbau principle: Electrons fill the lowest-energy molecular orbitals first.
- Pauli exclusion principle: Each molecular orbital can hold a maximum of two electrons with opposite spins.
- Hund's rule: When filling degenerate (equal-energy) molecular orbitals, electrons occupy them singly with parallel spins before pairing up.
The same quantum mechanical principles that govern electron configuration in atoms apply directly to molecules. This consistency is one of the strengths of MO theory.
Key Properties Summarised
The theory can be condensed into seven essential properties:
(I) Electrons in a molecule occupy molecular orbitals, just as electrons in an atom occupy atomic orbitals.
(II) Atomic orbitals of comparable energies and proper symmetry combine to form molecular orbitals.
(III) An atomic orbital is monocentric (influenced by one nucleus), while a molecular orbital is polycentric (influenced by two or more nuclei).
(IV) The number of molecular orbitals formed equals the number of combining atomic orbitals. When two atomic orbitals combine, two molecular orbitals result — one bonding and one antibonding. …